Periodic Table: Patterns

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Periodic Table Patterns
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John Newlands

• 1864
• arranged elements in octaves
• worked for some elements, but not all

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Dimitri Mendeleev

• Julius Lothar Meyer
• both independently created a form of the periodic
  table

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Dimitri Mendeleev

• Mendeleev usually credited because he showed how
  useful the table could be to predict properties

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Using the periodic table

• The periodic table has so much information
  contained in its organization that it will become
  your most valuable resource!!!

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Periodic Law

• Elements are arranged by atomic number, there is
  a periodic repetition in their physical and
  chemical properties.

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Periodic Table

• Horizontal rows
• Called periods
• 7 periods
• Show the number of the shell

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Periodic Table

• Vertical columns
• Called groups
• Show the valance electrons
• Elements in a group have similar physical and
  chemical properties

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Periodic Trends

• We have already learned one periodic trend-
• the way electrons are organized in atoms

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Some exceptions

• Draw the electron configuration of
• Cr
• Cu

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Single electron

• The energy of a single electron reflects how
  tightly bound that e- is to the nucleus (more
  negative energy)

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Effective Charge

• We assume that the electrons are bound by a
  positive charge Z
• But there is an effective charge that e- feels

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Effective Charge

• The effective charge (Zeff) represents the net
  effect of the attraction of the nuclear charge
  and the repulsion of the other e-s

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Effective Charge

• An electron closer to the nucleus would feel more
  of the positive charge

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Shielding

• The protection by the inner electrons so that
  the outer electrons do not feel the full nuclear
  charge

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Orbital Filling

• Shielding helps explain why the orbitals are
  filled in the order that they are
• The lower energy orbitals (ones closer to the
  nucleus) are always filled first

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Orbital Shielding

• Orbitals to the inside of the other orbitals do a
  good job of shielding the outer electrons

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Polyelectronic Model

• The energy required to remove an electron from an
  atom depends on two factors

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Two Factors

• The effective nuclear charge
• The average distance of the electron from the
  nucleus

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Atomic radius Half the distance between the
nuclei of two atoms. (Atoms are the same and
bonded together)
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Element Characteristics
Radius size increases
Radius size decreases
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Element Characteristics

• Ionization energy
• Energy required to remove an electron from an
  atom
• Easier to remove an electron from group 1 than
  group 8
• Easier to remove electrons that are farther away
  from the nucleus (elements at the bottom of a
  group)

Ionization energy generally decreases
Ionization energy increases
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The general trend

• As we go across a period from left to right, the
  first ionization energy increases

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Why??

• Electrons added in the same principle quantum
  level do not completely shield the increasing
  nuclear charge

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Exceptions

• Decreases in ionization energy moving across
  (i.e. N to O) is because e- in 2s provide some
  shielding for the 2p

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Electron Affinity

• The energy change associated with the addition of
  an electron to a gaseous atom
• X(g) e- --gt X-(g)

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Sign for the energy

• Defined as energy change when electron is added
• if addition is exothermic then the sign is
  negative

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Electron affinity

• What is the trend going down a group?
• Generally becomes more positive b/c e- added at
  increasing distance

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Trends...

• Electron affinities generally become more
  negative from left to right across a period,
  there are several exceptions

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Element Characteristics

• Electronegativity
• It is a measure of the element's ability to
  attract the electrons which are in a bond

Electronegativity decreases
Electronegativity increases
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Element Characteristics
More nonmetallic
Liquids
Periodic Table
More metallic
Solids
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Periodic Table Trends Summary
Atomic radius decreases
Ionization energy and Electronegativity increase